1. Chemistry Chapter 5 Electrons in Atoms
The 1998 Nobel Prize in Physics was awarded "for the discovery of a new form of quantum fluid with fractionally charged excitations." At the left is a computer graphic of this kind of state.

2. The Puzzle of the Atom
Protons and electrons are attracted to each other because of opposite charges
Electrically charged particles moving in a curved path give off energy
Despite these facts, atoms don’t collapse

3. Wave-Particle Duality
JJ Thomson won the Nobel prize for describing the electron as a particle.
His son, George Thomson won the Nobel prize for describing the wave-like nature of the electron.
The electron is a particle!
The electron is an energy wave!

4. Confused??? You’ve Got Company!
“No familiar conceptions can be woven around the electron; something unknown is doing we don’t know what.”
Physicist Sir Arthur Eddington
The Nature of the Physical World
1934

5. The Wave-like Electron
The electron propagates through space as an energy wave. To understand the atom, one must understand the behavior of electromagnetic waves.
Electromagnetic Spectrum
Louis deBroglie

6. Electromagnetic radiation propagates through space as a wave moving at the speed of light.
C = speed of light, a constant (3.00 x 108 m/s)
 = frequency, in units of hertz (hz, sec-1)
 = wavelength, in meters

7. Long Wavelength = Low Frequency Low ENERGY Short Wavelength =
Wavelength Table
Short
Wavelength =
High Frequency
High ENERGY

8. Types of electromagnetic radiation:
Spectrum

9. Photoelectric Effect
Emission of electrons from a metal when light shines on the metal.
Must have a minimum frequency of light to excite the e-.
But why?????
Photoelectric effect

10. Ground state- lowest energy state of an atom
Excited state- atom has higher potential energy than the ground state.
Excited
state
Ground
state
Photoelectric effect

11. Particle description of light
1900 Max Planck- emission of light by a hot object. Not emitted continuously, but in portions called quanta (minimum amt. of energy lost or gained by an atom)

12. The energy (E ) of electromagnetic radiation is directly proportional to the frequency () of the radiation.
E = h
E = Energy, in units of Joules (kg·m2/s2)
h = Planck’s constant (6.626 x J·s)
 = frequency, in units of hertz (hz, sec-1)

13. 1905 Albert Einstein- duel nature, light is a stream of particles called photons.
Has 0 rest mass and carries a quanta of energy.
E photon = hν

14. EM radiation is only absorbed in whole numbers of photons
In order for an electron to be emitted from a metal, it must be struck by a photon of minimal frequency.

15. If the photons frequency is below the min
If the photons frequency is below the min. then the electron will remain bound to the metal.
Different metals require different frequencies.

16. Spectroscopic analysis of the visible spectrum…
…produces all of the colors in a continuous spectrum
Results when the gas pressures are higher. Generally, solids, liquids, or dense gases emit light at all wavelengths when heated.

17. Hydrogen spectrum… …produces a “bright line” or emission spectrum
 produced by thin gases in which the atoms do not experience many collisions (because of the low density).
the “fingerprint” of light

18. Absorption spectrum
Occurs when light passes through a cold, dilute gas. The re-emitted light is not emitted in the same direction as the absorbed photon, this gives rise to dark lines (absence of light) in the spectrum.

19. Why had the H2 atoms given off only specific frequencies of light?
Quantum theory- an H2 atom emits a photon of radiation when it falls back down.

20. Electron transitions involve jumps of definite amounts of energy.
This produces bands
of light with definite
wavelengths.

21. The Bohr Model of the Atom
I pictured electrons orbiting the nucleus much like planets orbiting the sun.
But I was wrong! They’re more like bees around a hive.
WRONG!!!
Neils Bohr
Could only explain hydrogen and could not explain
how atoms bonded….

22. Quantum Model of the atom
Since light can behave as a wave and a particle, can particles of matter behave as waves?
Sure, we can call them
Matter waves!!

23. Quantum Mechanics Classical mechanics describes the motions
of bodies much larger than atoms, while
quantum mechanics describes the motion of
Subatomic particles and atoms as waves
Dr. Quantum

24. Heisenberg Uncertainty Principle
“One cannot simultaneously determine both the position and momentum of an electron.”
You can find out where the electron is, but not where it is going.
OR…
You can find out where the electron is going, but not where it is!
Werner
Heisenberg
entanglement

25. An orbital is a region within an atom where there is a probability of finding an electron. This is a probability diagram for the s orbital in the first energy level…
Orbital shapes are defined as the surface that
contains 90% of the total electron probability.

26. Schrodinger Wave Equation
Equation for probability of a single electron being found along a single axis (x-axis)
Erwin Schrodinger
Cat in a box
Big Bang Theory

27. Quantum Numbers
Each electron in an atom has a unique set of 4 quantum numbers which describe it.
Principal quantum number
Angular momentum quantum number
Magnetic quantum number
Spin quantum number
1 minute physics Entanglement

28. Pauli Exclusion Principle
No two electrons in an atom can have the same four quantum numbers.
Wolfgang
Pauli

29. Principal Quantum Number
Generally symbolized by n, it denotes the shell (energy level) in which the electron is located.
Number of electrons that can fit in a shell:
2n2

30. Angular Momentum Quantum Number
The angular momentum quantum number, generally symbolized by l, denotes the orbital (subshell) in which the electron is located.

31. n= 1 s orbital (spherical shape)
n= s & p orbital (p is dumbbell shaped)
n= 3 s,p, & d orbital
n= s,p,d, & f orbital

32. Magnetic Quantum Number
The magnetic quantum number, generally symbolized by m, denotes the orientation of the electron’s orbital with respect to the three axes in space.

33. Only one s because it’s spherical
p- can be on 3 axis (x, y, z) px, py, pz
d- five different orbitals
f – seven different orbitals

35. Assigning the Numbers
The three quantum numbers (n, l, and m) are integers.
The principal quantum number (n) cannot be zero.
n must be 1, 2, 3, etc.
The angular momentum quantum number (l) can be any integer between 0 and n - 1.
For n = 3, l can be either 0, 1, or 2.
The magnetic quantum number (m) can be any integer between -l and +l.
For l = 2, m can be either -2, -1, 0, +1, or +2.

36. Principle, angular momentum, and magnetic quantum numbers: n, l, and ml

37. Spin Quantum Number
Spin quantum number denotes the behavior (direction of spin) of an electron within a magnetic field.
Possibilities for electron spin:

38. An orbital is a region within an atom where there is a probability of finding an electron. This is a probability diagram for the s orbital in the first energy level…
Orbital shapes are defined as the surface that
contains 90% of the total electron probability.

39. Schrodinger Wave Equation
Equation for probability of a single electron being found along a single axis (x-axis)
Erwin Schrodinger

40. Heisenberg Uncertainty Principle
“One cannot simultaneously determine both the position and momentum of an electron.”
You can find out where the electron is, but not where it is going.
OR…
You can find out where the electron is going, but not where it is!
Werner
Heisenberg

41. Sizes of s orbitals Orbitals of the same shape (s, for instance) grow
larger as n increases…
Nodes are regions of low probability within an
orbital.

42. Orbitals in outer energy levels DO penetrate into lower energy levels.
Penetration #1
This is a probability
Distribution for a
3s orbital.
What parts of the
diagram correspond
to “nodes” – regions
of zero probability?

43. Which of the orbital types in the 3rd energy level
Does not seem to have a “node”?
WHY NOT?
Penetration #2

44. The s orbital has a spherical shape centered around
the origin of the three axes in space.
s orbital shape

45. P orbital shape There are three dumbbell-shaped p orbitals in
each energy level above n = 1, each assigned to
its own axis (x, y and z) in space.

46. Things get a bit more complicated with the five d orbitals that are found in the d sublevels beginning with n = 3. To remember the shapes, think of “double dumbells”
d orbital shapes
…and a “dumbell
with a donut”!

47. Shape of f orbitals

48. Orbital filling table

49. Electron configuration of the elements of the first three series

50. Irregular confirmations of Cr and Cu
Chromium steals a 4s electron to half
fill its 3d sublevel
Copper steals a 4s electron to FILL
its 3d sublevel